So I had a test question the other day, and it essentially went as follows:

A student measures the concentration of a $\ce{HCl}$ solution to be $10^{-9}$M by using a pH meter. Is the meter wrong?

This means that the $\ce{HCl}$ solution has a pH of 9, which seems unreasonable, yet that is not what I said. I said that the meter is not wrong. My reasoning was that you could have an $\ce{HCl}$ solution so incredibly dilute that the pH would in fact be 9. Since molarity is defined as moles per liter, if you hold the moles constant and increase the volume of solution, you will eventually reach a very tiny molarity.

In essence what I was wondering was: Is my reasoning correct?

Is this solution considered basic?

Can this be generalized such that any originally acidic solution is considered basic at a low enough concentration, or a high enough dilution?


2 Answers 2


Unfortunately, your reasoning is wrong, because you forgot to take into account the acidity of water. While in most cases we can ignore the acidity of water as the hydrogen ion concentration contribution by $\ce{HCl}$ dominates, at an extremely low concentration of $\ce{HCl},$ water becomes the main $\ce{H+}$ contributor.

It is well known that water has a $\pu{10^{-7} M}$ concentration of hydrogen ions at $\ce{25^\circ C}.$ Therefore, water's effect on the acidity of the solution is significant when looking at solutions of $\pu{10^{-6} M}$ or lower concentrations of strong acids. For example, if there is $\pu{10^{-9} M } \ce{HCl,}$ then the total $\ce{H+}$ concentration would still nearly be $\pu{10^{-7} M}.$ Therefore, no matter how much you dilute the acid, it can never turn into a basic solution.

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    $\begingroup$ Your answer is correct. At the same time, I think the OP's reasoning is also correct. You can dilute HCl to whatever concentration you want and therefore the meter is not wrong. What is wrong wtih that reasoning? $\endgroup$ Commented May 29, 2017 at 12:10
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    $\begingroup$ The concentration could be anything but it is impossible to measure an HCl concentration of 10^-9 M using a pH meter. So the meter is wrong. $\endgroup$
    – Paul
    Commented May 29, 2017 at 12:15
  • $\begingroup$ @Paul - You mean, assuming that the HCl is being diluted in water, the meter must be wrong. Perhaps some other solvent could be used to dilute the acid? At a minimum, using D2O should allow you to approach a pH of ~7.5. $\endgroup$
    – aroth
    Commented May 29, 2017 at 14:57
  • $\begingroup$ @aroth It won't, because D2O won't produce hydrogen ions when it dissociates. $\endgroup$ Commented May 29, 2017 at 22:03

Assuming this is an aqueous solution, you cannot have an HCl sample with pH 9. Consider: what is the pH of pure water? Of course, it is pH 7. As you add more and more solvent to your HCl solution, as you say, it becomes more dilute, but the autoionization of water eventually dominates and the pH will asymptote to 7 (from below since you started as an acid). The pH of a sample will always be between the pH's of the pure constituents. A solution with a pH greater than 7 is basic.

Now, let's say you had a sample of DCl (that is, deuterated HCl). If you could monitored the pD (analogous to pH) it could eventually get to 9, but the total acidity of the solution p(D+H) would go to 7, and it would always be acidic or neutral.


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